Physical Science 21
Laboratory
Lab 5: Spectra, Fingerprinting the Elements
Before coming to lab: Read Chapter 2.1 to 2.9 (on pages 30-36) and particularly Chapter
3.2 to 3.5 (on pages 63-67) of Astronomy (6th ed.) by Dinah Moché.
Objectives: This lab will introduce spectra, an important tool of astronomy. It will show
that each element has a characteristic spectrum, which is unique to that element.
Astronomers can therefore use spectra to tell what stars are made of—without having to
travel to the stars. This lab will also show the difference between continuous, emission,
and absorption spectra, which are made by planets, nebulae, and stars, respectively.
Introduction: When white light passes through a prism, it breaks down into its
component colors. A rainbow is an example of this in nature: after a rainstorm, the air is
still full of drops of water, which break sunlight down into its colors. This band of colors
is called a spectrum. The plural of “spectrum” is “spectra.” Using spectra to learn about
light is called spectroscopy.
Gustav Kirchhoff and Robert Bunsen were two German chemists. (Bunsen invented the
Bunsen burner, the gas burner still used in chemistry labs today.) In 1860, they found
that different hot gases gave off different colors of light, which they could distinguish
clearly in spectra. Kirchhoff summarized their findings into Kirchhoff’s laws:
(1) Hot gas that is transparent, such as a nebula, gives off light of only certain colors.
Often the gas emits only thin bands, or lines, of colors. This pattern of colors is called an
emission spectrum, or bright-line spectrum. Gases of different chemical composition
make different patterns of lines (emission spectra).
(2) A hot object that is opaque (not transparent to light), such as a planet, most solids,
liquids, and dense gases, gives off a continuous spectrum. A continuous spectrum is just a
continuous band of color with no lines, such as a rainbow.
(3) When a continuous spectrum shines through a transparent gas, dark lines appear in the
spectrum. This is an absorption spectrum, or dark-line spectrum. An absorption spectrum
can be thought of as a continuous spectrum, minus an emission spectrum. A star is made
of gas that is held together by its own gravity. Because a star is dense and opaque inside,
but is transparent outside, stars have absorption spectra.
Lab 5: Spectra, Fingerprinting the Elements
Light has properties of both waves and particles. So do atoms and electrons: the
realization that everything smaller than a molecule has properties of both waves and
particles is a central finding of the modern science of quantum mechanics. In this
experiment we will use both the wave and particle properties of light:
(1) Wavelength is the distance from one crest of a wave to the next crest. The symbol for
wavelength is λ.
(2) The color of light depends on its wavelength. Visible light has wavelengths between
400 nanometers (for violet light) to 780 nanometers (for the deepest red the eye can see.)
(1 nanometer = 1 nm = 1 billionth of a meter = 1 x 10-9 m.) White light is all the colors,
mixed together. A spectrum made from white light has the colors red, orange, yellow,
green, blue, and violet. (To remember the colors in order, remember ROY G. BV.)
(3) Particles of light are called photons. The energy, E, of a photon is related to its
wavelength, λ, such that:
E = hc/λ
where h is Planck’s constant and c is the speed of light in a vacuum. Notice that there is
an inverse relation between energy and wavelength: high-energy photons have short
wavelengths, and low-energy photons have long wavelengths. In other words, photons of
violet light have higher energy than photons of red light.
To understand how we can “fingerprint” the elements, we need to know how atoms emit
light. In 1897, J. J. Thompson showed that atoms are not indivisible, as was previously
thought. Thompson found that atoms have inside them even smaller particles, which are
now called electrons. In 1911, Ernest Rutherford found that nearly all the mass in an
atom is concentrated in its center, which is called the nucleus. In 1913, Niels Bohr found
that the electrons in an atom orbit the nucleus, in a way that reminded him of how the
planets in the Solar System orbit the Sun. However, Bohr soon realized that the electrons
can travel only in certain allowed orbits. This was the beginning of quantum mechanics:
“quantum” means “an allowed amount.”
Each of these orbits has an energy associated with it. An electron that orbits close to the
nucleus has a low energy, and an electron orbiting far from the nucleus has a higher
energy. If an electron jumps from a high-energy orbit to a low-energy orbit, it gives up its
extra energy in the form of a single photon. This process is called emission, since the
atom emits, or loses, energy. The energy carried away by the photon matches the
difference in the energy between the two orbits. This photon has a wavelength, λ,
corresponding to its energy E, such that E = hc/λ, as we’ve seen above.
The exact arrangement of the orbits and the energy differences between them differs,
between atoms of different chemical elements. Each element therefore has a unique set
of photons that it can emit. This gives each element its own spectrum: different elements
show different patterns of lines in their spectra. Spectra can therefore reveal the chemical
composition of things that make light, such as stars and nebulae.
Lab 5: Spectra, Fingerprinting the Elements
Glass can bend light. The amount of bending, or refraction, depends on the wavelength
of the light. This is how a prism separates white light into colors.
White Light
Red Light
Prism
Violet light
Figure 1. A glass prism disperses white light into a continuous spectrum.
A diffraction grating can also separate these colors. Gratings are easier to use than
prisms, since they can be made into eyeglasses.
Procedure
Use your grating to look at an incandescent light bulb (such as in an overhead projector).
Which colors are bent most? _______________ Which are bent least? _______________
(Notice that the above diagram, while illustrative, won’t necessarily work in every detail
here, since we’re using a grating, not a prism. This is to encourage you to try things out
yourself, and to rely on your own observations!)
You will see either a continuous or a line spectrum. Which do you see? ______________
Now let’s do some real spectroscopic identification of elements. In the lab, there should
be several lamps of hot, transparent gas. Look at these sources with your gratings
carefully. Make drawings of the lines on the form on the next page, and to try to maintain
the relative separation of these lines. Most of the lamps will have labels on them that tell
which gases they contain. One or two will not: use your drawings to identify the gases
that made those spectra.
Mystery Lamp 1: ____________________ Mystery Lamp 2: ____________________
Are these emission spectra or absorption spectra? _______________________________
QUESTION: Hot opaque gas, such as inside a star, will produce a continuous spectrum.
This light then travels through the transparent atmosphere of the star before it reaches us.
Do you expect stars to have emission spectra or absorption spectra? _________________
QUESTION: Look out the window. Do the pink (or white) street lights (from mercury
vapor lamps) have absorption or emission spectra? ______________________________
QUESTION: Look out the window. Do the yellow street lights (from sodium vapor
lamps) have absorption or emission spectra? ____________________________________
Lab 5: Spectra, Fingerprinting the Elements
Lab 5: Spectra, Fingerprinting the Elements